Enthalpy Terms - Concept Map

Surroundings:

Refers to everything else in the universe

Extra Terms

• Latent heat
• Latent heat is the heat energy absorbed or released during a phase change of a substance without a change in temperature, such as during melting or vaporization
• Neutralization
• Neutralization is a chemical reaction between an acid and a base, resulting in the formation of water and a salt, and leading to the neutralization of acidic and basic properties
• The Chapman Cycle
• The Chapman Cycle is a process occurring in the stratosphere where ozone (O3) is formed and destroyed through a series of reactions involving ultraviolet (UV) radiation and oxygen molecules.

• Pressure
• Pressure is the force exerted per unit area, resulting from the collisions of gas molecules with the walls of their container
• Equilibrium
• Equilibrium refers to a state in which opposing forces or processes are balanced, resulting in a stable system with no net change over time
• First Ionization Energy
• The first ionization energy is the energy required to remove the outermost electron from a neutral atom in its gaseous state, forming a positively charged ion
• Electron Affinity
• Electron affinity is the energy change that occurs when a neutral atom in the gaseous phase gains an electron to form a negatively charged ion.

Energy is the ability to do work.

Chemical Energy

• Work
• Work is defined as the transfer of energy that occurs when a force is applied to an object, causing it to move in the direction of the force.
• Heat is a mode of energy transfer as a result of temperature differences
• Produces an increase in average kinetic energy of particles
• Causes an increase in disorder of particles
• Chemical Energy

• Energy absorbed/released during chemical reactions
• Energy absorbed is endothermic
• Energy released is exothermic

System vs Surroundings

Exothermic vs Endothermic

System

• Refers to specifically the area of interest
• Open Systems allow exchange of matter and energy.
• Closed Systems allow exchange of energy only.

Exothermic

• Energy released is exothermic
• Making bonds is exothermic
• Lattice enthalpies break apart a solid ionic lattice into its gaseous ions, which is an endothermic process

Endothermic

• Energy absorbed is endothermic
• Breaking bonds is endothermic
• Gaseous ions coming together to form an ionic lattice is an exothermic process

Enthalpy of Formation

Standard Enthalpy Changes

• ΔH⦵ is used to describe enthalpy changes in chemical reactions
• 100kPa
• 1 mol dm-3 concentration for all solutions
• All substances in their standard states
• Standard States is the pure form of a substance under standard conditions of 298 K (25 °C) and 100 kPa

• Standard Enthalpy of Formation (ΔHf⦵): the enthalpy change that occurs when one mole of the substance is formed from its elements in their standard states.
• Gives a measure of the stability of a substance relative to its elements
• Can be used to calculate the enthalpy change of all reactions (hypothetical or real)

Enthalpy

ΔG=ΔH−TΔS

Average Bond enthalpies

Hess’s Law

• ΔH⦵rxn = ∑ΔH⦵(bonds broken)-∑ΔH⦵(bonds formed)
• Bond enthalpies are the energy needed to break one mole of bonds in gaseous molecules under standard conditions averaged over similar compounds

• Sometimes enthalpy changes of a reaction are not measured directly.
• The ΔH of a reaction is calculated from the known enthalpy changes of other reactions.
• Regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes.

• q
• Enthalpy Change
• H
• Enthalpy of a system
• Standard Enthalpy Changes
• Enthalpy of Formation
• Average Bond enthalpies
• Hess’s Law
• Enthalpies of Solution (ΔHsol⦵)
• Hydration Enthalpy (ΔHhyd⦵)
• Lattice Enthalpy (ΔHlat⦵)

Enthalpies of Solution, Hydration Enthalpy, Lattice Enthalpy

Born-Haber Cycle

• The Born-Haber cycle is a method to determine lattice enthalpies by accounting for the energy changes during the formation of an ionic compound from its elements in their standard states.

• Enthalpies of Solution (ΔHsol⦵)
• Enthalpy of solution is the heat energy absorbed or released when a solute dissolves in a solvent under standard conditions
• Hydration Enthalpy (ΔHhyd⦵)
• Hydration enthalpy is the heat energy absorbed or released when ions are hydrated by solvent molecules under standard conditions
• Lattice Enthalpy (ΔHlat⦵)
• Lattice enthalpy is the energy released when gaseous ions combine to form a solid ionic lattice under standard conditions.

• ΔG = Gibbs Free Energy
• Gibbs free energy is a thermodynamic function that represents the maximum reversible work that can be obtained from a system at constant temperature and pressure.
• ΔH = Enthalpy of a system
• ΔS = Entropy
• T = temperature

Entropy

• Entropy is a measure of the disorder or randomness of a system, representing the number of ways the system's particles can be arranged.

Enthalpy Terms

by Maxwell Carter

q=mcΔt

• q = Enthalpy change
• The amount of heat transferred
• c = Specific heat
• It quantifies the amount of heat energy required to change the temperature of a unit mass of the substance by one degree Celsius.
• t = Temperature
• The average kinetic energy of the particles in the system.